Chapter 17

Concept Review What is meant by a half-reaction?

Why can't a wire perform the same function as a porous separator in an electrochemical cell?

In a voltaic cell, why is the cathode labeled the positive terminal and the anode the negative terminal?

Balance the following half-reactions by adding the appropriate number of electrons. Identify the oxidation half-reactions and the reduction half- reactions. a. \(\quad \mathrm{Br}_{2}(\ell) \rightarrow 2 \mathrm{Br}^{-}(a q)\) b. \(\mathrm{Pb}(s)+2 \mathrm{Cl}^{-}(a q) \rightarrow \mathrm{PbCl}_{2}(s)\) c. \(\mathrm{O}_{3}(g)+2 \mathrm{H}^{+}(a q) \rightarrow \mathrm{O}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(\ell)\) d. \(\mathrm{H}_{2} \mathrm{S}(g) \rightarrow \mathrm{S}(s)+2 \mathrm{H}^{+}(a q)\)

Balance the following half-reactions by adding the appropriate number of electrons. Which are oxidation half-reactions and which are reduction half- reactions? a. \(\mathrm{TiO}^{2+}(a q)+2 \mathrm{H}^{+}(a q) \rightarrow \mathrm{Ti}^{3+}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell)\) b. \(\mathrm{S}_{4}^{2+}(a q) \rightarrow 2 \mathrm{S}_{2}^{2-}(a q)\) c. \(\mathrm{VO}_{2}{\underline{\phantom{xx}}}^{+}(a q)+2 \mathrm{H}^{+}(a q) \rightarrow \mathrm{VO}^{2+}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell)\) d. \(\mathrm{Fe}(\mathrm{CN})_{6}^{3-} \rightarrow \mathrm{Fe}(\mathrm{CN})_{6}{\underline{\phantom{xx}}}^{4-}(a q)\)

Write a half-reaction for the oxidation of magnetite \(\left(\mathrm{Fe}_{3} \mathrm{O}_{4}\right)\) to hematite \(\left(\mathrm{Fe}_{2} \mathrm{O}_{3}\right)\) in acidic groundwater.

Write a half-reaction for the oxidation of the manganese in \(\mathrm{MnCO}_{3}\) to \(\mathrm{MnO}_{2}\) in neutral groundwater, where the principal carbonate species is \(\mathrm{HCO}_{3}^{-}.\)

Lithium-Iron Sulfide Batteries Voltaic cells with nonaqueous electrolytes and based on the oxidation of \(\mathrm{Li}(s)\) to \(\mathrm{Li}_{2} \mathrm{S}(s)\) and the reduction of \(\mathrm{FeS}_{2}(s)\) to \(\mathrm{Fe}(s)\) and \(\mathrm{S}^{2-}\) ions provide the same voltage as traditional alkaline batteries but offer more storage capacity. a. Write half-reactions for the cell's anode and cathode. b. Write a balanced cell reaction. c. Diagram the cell.

In a flow battery, two solutions are pumped into each voltaic cell, and oxidation and reduction half-reactions take place at two inert electrodes. One such cell is based on the reduction of \(\mathrm{Sn}^{4+}(a q)\) to \(\mathrm{Sn}^{2+}(a q)\) and the oxidation of \(\mathrm{Fe}^{2+}(a q)\) to \(\mathrm{Fe}^{3+}(a q)\) a. Write half-reactions for the cell's anode and cathode. b. Write a balanced cell reaction. c. How many electrons are transferred in the cell reaction?

Zinc-Nickel Batteries Replacing cadmium with zinc in a NiCad battery avoids the use of the toxic element cadmium. A voltaic cell based on the reaction between \(\mathrm{NiO}(\mathrm{OH})(s)\) and \(\mathrm{Zn}(s)\) in alkaline electrolyte \(\left(\mathrm{OH}^{-}\right)\) produces \(\mathrm{Ni}(\mathrm{OH})_{2}(s)\) and \(\mathrm{Zn}(\mathrm{OH})_{2}(s).\) a. Write half-reactions for the anode and cathode. b. Write a balanced cell reaction. c. Diagram the cell.

Super Iron Batteries In \(1999,\) scientists in Israel developed a battery based on the following cell reaction with iron(V1), nicknamed "super iron": \(2 \mathrm{K}_{2} \mathrm{FeO}_{4}(a q)+3 \mathrm{Zn}(s) \rightarrow \mathrm{Fc}_{2} \mathrm{O}_{3}(s)+\mathrm{ZnO}(s)+2 \mathrm{K}_{2} \mathrm{ZnO}_{2}(a q).\) a. Determine the number of electrons transferred in the cell reaction. b. What are the oxidation states of the transition metals in the reaction? c. Draw the cell.

Why is \(\mathrm{O}_{2}\) a stronger oxidizing agent in acid than in base? Use standard reduction potentials from Appendix 6 to support your answer.

From the table of standard reduction potentials in Appendix 6 a. Select an oxidizing agent that will oxidize \(\mathrm{Cr}(s)\) to \(\mathrm{Cr}^{3+}(a q)\) but not \(\mathrm{Cd}(s)\) to \(\mathrm{Cd}^{2+}(a q)\) b. Select a reducing agent that will reduce Br \(_{2}(\ell)\) to \(\mathrm{Br}^{-}(a q)\) but not \(\mathrm{I}_{2}(s)\) to \(1^{-}(a q).\)

If a piece of silver is placed in a solution in which \(\left[\mathrm{Ag}^{+}\right]\) \(=\left[\mathrm{Cu}^{2+}\right]=1.00 \mathrm{M},\) will the following reaction procecd spontancously? $$ 2 \Lambda g(s)+\mathrm{Cu}^{2+}(a q) \rightarrow 2 \mathrm{Ag}^{+}(a q)+\mathrm{Cu}(s) $$

Voltaic cells based on the following pairs of half-reactions are prepared so that all reactants and products are in their standard states. For each pair, write a balanced equation for the cell reaction, and identify which half- reaction takes place at the anode and which at the cathode. a. \(\mathrm{Hg}^{2+}(a q)+2 \mathrm{c}^{-} \rightarrow \mathrm{Hg}(\ell)\) \(\mathrm{Zn}^{2+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{Zn}(s)\) b. \(\mathrm{ZnO}(s)+\mathrm{H}_{2} \mathrm{O}(\ell)+2 \mathrm{e}^{-} \rightarrow \mathrm{Zn}(s)+2 \mathrm{OH}^{-}(a q)\) \(\mathrm{Ag}_{2} \mathrm{O}(s)+\mathrm{H}_{2} \mathrm{O}(\ell)+2 \mathrm{c}^{-} \rightarrow 2 \mathrm{Ag}(s)+2 \mathrm{OH}^{-}(a q)\) c. \(\mathrm{Ni}(\mathrm{OH})_{2}(s)+2 \mathrm{e}^{-} \rightarrow \mathrm{Ni}(s)+2 \mathrm{OH}^{-}(a q)\) \(\mathrm{O}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(\ell)+4 \mathrm{e}^{-} \rightarrow 4 \mathrm{OH}^{-}(a q)\)

Voltaic cells based on the following pairs of half-reactions are constructed. For each pair, write a balanced equation for the cell reaction, and identify which half-reaction takes place at each anode and cathode. a. \(\mathrm{Cd}^{2+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{Cd}(s)\) \(\mathrm{Ag}^{+}(a q)+\mathrm{c}^{-} \rightarrow \mathrm{Ag}(s)\) b. \(\mathrm{AgBr}(s)+\mathrm{c}^{-} \rightarrow \mathrm{Ag}(s)+\mathrm{Br}^{-}(a q)\) \(\mathrm{MnO}_{2}(s)+4 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{Mn}^{2+}(a q)+2 \mathrm{H}_{2} \mathrm{O}(\ell)\) c. \(\operatorname{Pt} C l_{4}^{2-}(a q)+2 e^{-} \rightarrow \operatorname{Pt}(s)+4 C l^{-}(a q)\) \(\mathrm{AgCl}(s)+\mathrm{c}^{-} \rightarrow \mathrm{Ag (s)+\mathrm{Cl}^{-}(a q)\)

The negative sign in Equation \(17.3\left(w_{\text {clec }}=-Q E_{\text {cell }}\right)\) seems to indicate that a voltaic cell with a positive cell potential does negative electrical work. How is this possible?

Starting with the appropriate standard free energies of formation from Appendix \(4,\) calculate the valucs of \(\Delta G^{\circ}\) and \(E_{\text {cell }}\) of the following reactions: a. \(2 \mathrm{Cu}^{+}(a q) \rightarrow \mathrm{Cu}^{2+}(a q)+\mathrm{Cu}(s)\) b. \(\mathrm{Ag}(s)+\mathrm{Fe}^{3+}(a q) \rightarrow \mathrm{Ag}^{+}(a q)+\mathrm{Fe}^{2+}(a q)\)

Starting with the appropriate standard free encrigics of formation from Appendix 4, calculate the values of \(\Delta G^{*}\) and \(E_{\text {cell of the following reactions: }}\) a. \(\mathrm{FeO}(s)+\mathrm{H}_{2}(g) \rightarrow \mathrm{Fe}(s)+\mathrm{H}_{2} \mathrm{O}(\ell)\) b. \(2 \mathrm{Pb}(s)+\mathrm{O}_{2}(g)+2 \mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow\) \(2 \mathrm{PbSO}_{4}(s)+2 \mathrm{H}_{2} \mathrm{O}(\ell)\)

Laptops The first generation of laptop computers was powered by nickcl-cadmium (NiCad) batterics, which generated \(1.20 \mathrm{V}\) based on the following cell reaction: \(\mathrm{Cd}(s)+2 \mathrm{NiO}(\mathrm{OH})(s)+2 \mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow\) $$ \mathrm{Cd}(\mathrm{OH})_{2}(s)+2 \mathrm{Ni}(\mathrm{OH})_{2}(s) $$ What is the value of \(\Delta G_{\text {cell }} ?\)

Nickel-Sodium Batteries Researchers in England are developing a battery for electric vehicles based on the reaction between \(\mathrm{NiCl}_{2}(s)\) and \(\mathrm{Na}(\mathrm{s}):\) $$ 2 \mathrm{Na}(s)+\mathrm{NiCl}_{2}(s) \rightarrow \mathrm{Ni}(s)+2 \mathrm{NaCl}(s) $$ The cells in the battery produce \(2.58 \mathrm{V}\) a. Assign oxidation numbers to each element in the nickel and sodium compounds. b. How many electrons are transferred in the overall reaction? c. What is the value of \(\Delta G_{\text {cen }} ?\)

Thomas Edison, the inventor of the incandescent lightbulb, developed a voltaic cell that delivers \(1.4 \mathrm{V}\) of cell potential in an alkaline electrolyte based on the following cell reaction: \(3 \mathrm{Fe}(s)+8 \mathrm{NiO}(\mathrm{OH})(s)+4 \mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow\) $$ 8 \mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{Fe}_{3} \mathrm{O}_{4}(s) $$ a. Assign oxidation numbers to each element in each of the nickel and iron compounds. b. How many electrons are transferred in the overall reaction? c. What is the value of \(\Delta G_{\text {cell }} ?\)

Starting with standard potentials listed in Appendix 6 calculate the values of \(E_{\text {cell }}^{\circ}\) and \(\Delta G^{\circ}\) of the following reactions. a. \(\mathrm{Cu}(s)+\mathrm{Sn}^{2+}(a q) \rightarrow \mathrm{Cu}^{2+}(a q)+\mathrm{Sn}(s)\) b. \(\mathrm{Zn}(s)+\mathrm{Ni}^{2+}(a q) \rightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Ni}(s)\)

A Reference Point: The Standard Hydrogen Electrode What is the function of platinum in the standard hydrogen electrode?

The Effect of Concentration on \(E_{\mathrm{cell}}\) Why does the operating cell potential of most batteries change little until the battery is nearly discharged?

The standard potential of the \(\mathrm{Zn} / \mathrm{Cu}^{2+}\) cell reaction $$ \mathrm{Zn}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Cu}(s) $$ is \(1.10 \mathrm{V}\). Would the potential of the cell differ from \(1.10 \mathrm{V}\) if the concentrations of both \(\mathrm{Cu}^{2+}\) and \(\mathrm{Zn}^{2+}\) were \(0.25 M ?\)

Using the appropriate standard potentials from Appendix 6, determine the equilibrium constant for the following reaction at \(298 \mathrm{K}:\) $$ \mathrm{Fe}^{3+}(a q)+\mathrm{Cr}^{2+}(a q) \rightarrow \mathrm{Fe}^{2+}(a q)+\mathrm{Cr}^{3+}(a q) $$

Using the appropriate standard potentials from Appendix \(6,\) determine the equilibrium constant at \(298 \mathrm{K}\) for the following reaction between \(\mathrm{Mn} \mathrm{O}_{2}\) and \(\mathrm{Fe}^{2+}\) in acid solution: \(4 \mathrm{H}^{+}(a q)+\mathrm{MnO}_{2}(s)+2 \mathrm{Fe}^{2+}(a q) \rightarrow\) $$ \mathrm{Mn}^{2+}(a q)+2 \mathrm{Fe}^{3+}(a q)+2 \mathrm{H}_{2} \mathrm{O}(\ell) $$

If the potential of a hydrogen electrode based on the half-reaction $$ 2 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2}(g) $$ is \(0.000 \mathrm{V}\) at \(\mathrm{pH}=0.00,\) what is the potential of the same electrode at \(\mathrm{pH}=7.00 ?\)

Glucose Metabolism The standard potentials for the reduction of nicotinamide adenine dinuclcotide (NAD") and oxaloacetate (reactants in the multistep metabolism of glucose) are as follows: $$ \mathrm{NAD}^{+}(a q)+2 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{NADH}(a q)+\mathrm{H}^{+}(a q) $$ Oxaloacetate \((a q)+2 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \rightarrow\) malate \((a q)\) \(E^{*}=-0.166 \mathrm{V}\) a. Calculate the standard potential for the following reaction: Oxaloacetate \((a q)+\mathrm{NADH}(a q)+\mathrm{H}^{+}(a q) \rightarrow\) malate( \(a q)+\mathrm{NAD}^{+}(a q)\) b. Calculate the equilibrium constant for the reaction at \(298 \mathrm{K}.\)

A copper penny dropped into a solution of nitric acid produces a mixture of nitrogen oxides. The following reaction describes the formation of \(\mathrm{NO},\) one of the products: \(3 \mathrm{Cu}(s)+8 \mathrm{H}^{+}(a q)+2 \mathrm{NO}_{3}^{-}(a q) \rightarrow\) \(2 \mathrm{NO}(g)+3 \mathrm{Cu}^{2+}(a q)+4 \mathrm{H}_{2} \mathrm{O}(\ell)\) a. Starting with the appropriate standard potentials from Appendix \(6,\) calculate \(E_{\text {ren }}^{\circ}\) for this reaction. b. Calculate \(E_{\text {rxn }}\) at \(298 \mathrm{K}\) when \(\left[\mathrm{H}^{+}\right]=0.100 \mathrm{M}\) \(\left[\mathrm{NO}_{3}^{-}\right]=0.0250 \mathrm{M},\left[\mathrm{Cu}^{2+}\right]=0.0375 M,\) and the partial pressure of \(\mathrm{NO}=0.00150\) atm.

Chlorine dioxide \(\left(\mathrm{ClO}_{2}\right)\) is produced by the following reaction of chlorate \(\left(\mathrm{ClO}_{3}^{-}\right)\) with \(\mathrm{Cl}^{-}\) in acid solution: \(2 \mathrm{ClO}_{3}^{-}(a q)+2 \mathrm{Cl}^{-}(a q)+4 \mathrm{H}^{+}(a q) \rightarrow\) $$ 2 \mathrm{ClO}_{2}(g)+\mathrm{Cl}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) $$ a. Determine \(E^{\circ}\) for the reaction. b. The reaction at \(298 \mathrm{K}\) produces a mixture of gases in the reaction vessel in which \(P_{\mathrm{ClO}_{2}}=2.0\) atm and \(P_{\mathrm{C}_{1}}=1.00\) atm. Calculate \(\left[\mathrm{ClO}_{3}^{-}\right]\) if, at equilibrium, \(\left[\mathrm{H}^{+}\right]=\left[\mathrm{Cl}^{-}\right]=10.0 \mathrm{M}.\)

The oxidation of \(\mathrm{NH}_{4}^{+}\) to \(\mathrm{NO}_{3}^{-}\) in acid solution is described by the following equation: \(\mathrm{NH}_{4}^{+}(a q)+2 \mathrm{O}_{2}(g) \rightarrow \mathrm{NO}_{3}^{-}(a q)+2 \mathrm{H}^{+}(a q)+\mathrm{H}_{2} \mathrm{O}(\ell)\) a. Calculate \(E^{\circ}\) for the overall reaction. b. If the reaction is in equilibrium with air \(\left(P_{\mathrm{O}_{1}}=0.21 \mathrm{atm}\right)\) at pH \(5.60,\) what is the ratio of \(\left[\mathrm{NO}_{3}^{-}\right]\) to \(\left[\mathrm{NH}_{4}^{+}\right]\) at \(298 \mathrm{K} ?\)

What is the value of \(E^{\circ}\) for the following reaction? $$ 2 \mathrm{AgCl}(s)+\mathrm{H}_{2}(g) \rightarrow 2 \mathrm{Ag}(s)+2 \mathrm{HCl}(a q) $$

Relating Battery Capacity to Quantities of Reactants One 12 -volt lead-acid battery has a higher ampere \cdot hour rating than another. Which of the following parameters are likely to be different for the two batteries? a. Individual cell potentials b. Anode half-reactions c. Total masses of electrode materials d. Number of cells e. Electrolyte composition f. Combined surface areas of their electrodes

In a voltaic cell based on the \(\mathrm{Zn} / \mathrm{Cu}^{2+}\) cell reaction, $$ \mathrm{Zn}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow \mathrm{Cu}(s)+\mathrm{Zn}^{2+}(a q) $$ there is exactly one mole of each reactant and product. A second cell based on the following cell reaction: $$ \mathrm{Cd}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow \mathrm{Cu}(s)+\mathrm{Cd}^{2+}(a q) $$ also has exactly one mole of each reactant and product. Which of the following statements about these two cells is true? a. Their cell potentials are the same. b. The masses of their electrodes are the same. c. The quantities of electric charge that they can produce are the same. d. The quantities of electric energy that they can produce are the same.

Which of the following voltaic cells, \(\boldsymbol{A}\) or \(\mathbf{B},\) will produce the greater quantity of electric charge per gram of anode material? Cell \(\mathbf{A}: \mathbf{C d}(s)+2 \mathrm{NiO}(\mathrm{OH})(s)+2 \mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow\) \(2 \mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{Cd}(\mathrm{OH})_{2}(\mathrm{s})\) Cell \(\mathbf{B}: 4 \mathrm{Al}(s)+3 \mathrm{O}_{2}(g)+6 \mathrm{H}_{2} \mathrm{O}(\ell)+4 \mathrm{OH}^{-}(a q) \rightarrow\) \(4 \mathrm{Al}(\mathrm{OH})_{4}-(a q)\)

Which of the following voltaic cell reactions, \(\mathbf{E}\) or \(\mathbf{F}\), delivers more electrical energy per gram of anode material at \(298 \mathrm{K} ?\) Reaction \(\mathbf{E}: \mathbf{Z n}(s)+2 \mathrm{NiO}(\mathrm{OH})(s)+2 \mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow\) $$ 2 \mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{Zn}(\mathrm{OH})_{2}(s) \quad E_{c c 11}^{*}=1.20 \mathrm{V} $$ Reaction \(\mathbf{F}: \mathrm{Li}(s)+\mathrm{MnO}_{2}(s) \rightarrow \mathrm{LiMnO}_{2}(s) \quad E_{\mathrm{call}}^{\circ}=3.15 \mathrm{V}\)

Which of the following voltaic cell reactions, \(\mathbf{G}\) or \(\mathbf{H}\), delivers more electrical energy per gram of anode material at \(298 \mathrm{K} ?\) Reaction \(\mathbf{G}: \mathbf{Z n}(\mathbf{s})+\mathbf{N i}(\mathbf{O H})_{2}(s) \rightarrow\) $$ \mathrm{Zn}(\mathrm{OH})_{2}(s)+\mathrm{Ni}(s) \quad E_{c \mathrm{ell}}^{*}=1.50 \mathrm{V} $$ Reaction \(\mathbf{H}: 2 \mathrm{Zn}(s)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{ZnO}(s) \quad E_{\text {cell }}^{*}=2.08 \mathrm{V}\)

Electrolytic Cells and Rechargeable Batteries The positive terminal of a voltaic cell is the cathode. However, the cathode of an electrolytic cell is connected to the negative terminal of a power supply. Explain this difference in polarity.

The anode in an electrochemical cell is defined as the electrode where oxidation takes place. Why is the anode in an electrolytic cell connected to the positive \((+)\) terminal of an external supply, whereas the anode in a voltaic cell battery is connected to the negative \((-)\) terminal?

The salts obtained from the evaporation of seawater can act as a source of halogens, principally \(\mathrm{Cl}_{2}\) and \(\mathrm{Br}_{2}\), through the electrolysis of the molten alkali metal halides. As the potential of the anode in an electrolytic cell is increased, which of these two halogens forms first?

Quantitative Analysis Electrolysis can be used to determine the concentration of \(\mathrm{Cu}^{2+}\) in a given volume of solution by electrolyzing the solution in a cell equipped with a platinum cathode. If all of the \(\mathrm{Cu}^{2+}\) is reduced to \(\mathrm{Cu}\) metal at the cathode, the increase in mass of the electrode provides a measure of the concentration of \(\mathrm{Cu}^{2+}\) in the original solution. To ensure the complete (99.9996) removal of the \(\mathrm{Cu}^{2+}\) from a solution in which \(\left[\mathrm{Cu}^{2+}\right]\) is initially about \(1.0 M,\) will the potential of the cathode (versus SHE) have to be more negative or less negative than \(0.34 \mathrm{V}\) (the standard potential for \(\mathrm{Cu}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Cu}\) )?

A quantity of electric charge deposits \(0.732 \mathrm{g}\) of \(\mathrm{Ag}(\mathrm{s})\) from an aqueous solution of silver nitrate. When that same quantity of charge is passed through a solution of a gold salt, \(0.446 \mathrm{g}\) of \(\mathrm{Au}(s)\) is formed. What is the oxidation state of the gold ion in the salt?

How long does it take to deposit a coating of gold \(1.00 \mu \mathrm{m}\) thick on a disk-shaped medallion \(4.0 \mathrm{cm}\) in diameter and \(2.0 \mathrm{mm}\) thick at a constant current of \(85 \mathrm{A} ?\) The density for the electroplating process is \(19.3 \mathrm{g} / \mathrm{cm}^{3} .\) The electroplating solution contains gold(111).

Oxygen Supply in Submarines Nuclear submarines can stay under water nearly indefinitely because they can produce their own oxygen by the electrolysis of water. a. How many liters of \(\mathrm{O}_{2}\) at \(298 \mathrm{K}\) and 1.00 bar are produced in \(1 \mathrm{h}\) in an electrolytic cell operating at a current of \(0.025 \mathrm{A} ?\) b. Could seawater be used as the source of oxygen in this electrolysis? Explain why or why not.

Calculate the minimum (least negative) cathode potential (versus SHE) needed to begin electroplating nickel from \(0.35 M \mathrm{Ni}^{2+}\) onto a piece of iron.

What is the minimum (least negative) cathode potential (versus SHE) needed to clectroplate silver onto cutlery in a solution of \(\mathrm{Ag}^{+}\) and \(\mathrm{NH}_{3}\) in which most of the silver ions are present as the complex \(\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2}^{+}\) and the concentration of \(\mathrm{Ag}^{+}(a q)\) is only \(3.50 \times 10^{-5} \mathrm{M} ?\)

Describe two advantages of hybrid (gasoline engine electric motor) power systems over all-electric systems based on fuel cells. Describe two disadvantages.

Describe three factors limiting the more widespread use of cars and other vehicles powered by fuel cells.